Post on 16-Oct-2021
Draft
Chemistry at Dalhousie circa 1868
Journal: Canadian Journal of Chemistry
Manuscript ID cjc-2017-0608.R1
Manuscript Type: Article
Date Submitted by the Author: 16-Nov-2017
Complete List of Authors: Grossert, J.; Dalhousie University White, Robert; Dalhousie University Ramaley, Louis; Dalhousie University Department of Chemistry
Is the invited manuscript for consideration in a Special
Issue?: Dalhousie
Keyword: chemistry at Dalhousie in 1868, early chemistry in Nova Scotia, George Lawson, organic chemicals from plants, common elements
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
1/27
Chemistry at Dalhousie circa 1868
J. Stuart Grossert, Robert L. White, and Louis Ramaley
Department of Chemistry
Dalhousie University
6274 Coburg Road, PO Box 15000
Halifax, NS, Canada B3H 4R2
Corresponding Author: J. Stuart Grossert
email: j.s.grossert@dal.ca
phone: 902 494 3305
FAX: 902 494 1310
Page 1 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
2/27
Abstract: This article describes details of classes at Dalhousie University in 1868‒1869, of the life
of George Lawson, the first Professor of Chemistry and Mineralogy, and of the wide range of
chemical concepts known at that time. A comprehensive set of lecture notes from Lawson's
chemistry course, written by a student, Alexander Russell, and held in the Dalhousie University
Archives, offers a wonderful insight into the state of chemical knowledge and how it was taught at
that time. Lawson began with general chemical principles followed by a detailed discussion of the
nonmetals. The second half of the class covered a range of metals followed by a small section on
mineralogy and a large section on organic and biological chemistry. Lawson used an older set of
atomic masses in which many, but not all, of the elements had masses one-half of the accepted
values today. When corrected for these errors, Lawson’s formulae, even for complex molecules
such as morphine, mostly agreed with contemporary usage. Examples of nomenclature, chemical
formulae, preparations, processes and properties are presented. A few examination questions are
given also. Even though the concepts involved in understanding chemical structure were just being
developed, the breadth and depth of descriptive chemical knowledge at that time was remarkable.
Key Words: chemistry at Dalhousie in 1868, early chemistry in Nova Scotia, George Lawson,
organic chemicals from plants, common elements
Page 2 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
3/27
Introduction
This issue of the Canadian Journal of Chemistry celebrates the 200th anniversary of the
founding of Dalhousie University. Contributions from Faculty and Alumni of the Department of
Chemistry highlight current aspects of research. Therefore, we felt it was appropriate to recognize
the long history of chemistry at Dalhousie by providing a perspective from one hundred and fifty
years ago. From 1838 on,1 chemistry was included in classes on Natural Philosophy while chemistry
as a subject has been taught since 1863. At that time, an infusion of cash allowed Dalhousie to
reorganize, hire new faculty, and expand its subject offerings. Chemistry was considered important
and was included in that expansion with the appointment of George Lawson as the first Professor of
Chemistry and Mineralogy. This paper presents a rare view through the eyes of a student of how
aspects of chemistry were taught in Canada a century and a half ago.
The 1868–69 Dalhousie Calendar lists only two degrees in the Faculty of Arts,2 a general BA
degree and an MA degree. The faculty and staff consisted of six professors (including the Principal
and Lawson), a tutor for Modern Languages and a janitor. Earning a BA degree required successful
completion of four winter terms which ran from late October until early April, with lectures in most
classes being held every day of the week. Required subjects included Classics, Mathematics,
Rhetoric, Logic and Psychology, Chemistry, Natural Philosophy (Physics), Modern Languages,
Metaphysics, Ethics, Political Economy and History, many of which were taught in more than one
year. Modern languages could be either German or French. Chemistry was taught in both the
second (junior chemistry) and third year (senior chemistry). By way of comparison, at this time
McGill University3 offered one Chemistry class of some 45 lectures to first-year BA students.
Queen's University4 offered one class to third-year BA students, whereas University College,
Toronto,5 offered several classes to both third- and fourth-year students.
In 1868–69 there were a total of 32 students enrolled at Dalhousie in the four years of the BA
degree plus a further 26 “general students” who were permitted to attend classes on payment of the
Page 3 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
4/27
requisite fee.2 In addition, the 22 students registered in the newly formed Faculty of Medicine were
required to take chemistry in their first year of study. The MA degree was awarded to individuals who
“maintained a good reputation” for three years after the BA and submitted an acceptable thesis,
accompanied by payment for the required diploma fee. According to Department of Chemistry
records, Dalhousie awarded the first Master of Arts Degree in Chemistry to James Forrest in 1871
and the initial First Class Honours Degree in Experimental Physics and Chemistry to Ebenezer
MacKay in 1886. It is interesting to note that in 1896 Ebenezer MacKay succeeded Lawson as the
second Professor of Chemistry at Dalhousie.
Professor George Lawson
Professor George Lawson was an interesting individual. Although formally appointed as
Professor of Chemistry and Mineralogy, he was a botanist at heart and published almost all of his
scientific work in botany.6 Born in Scotland in 1827, Lawson first pursued a career in the law, but
gave this up to study science at the University of Edinburgh in 1848.7-9 For the next decade he
continued his scientific studies there, acting as a demonstrator in the laboratory of John H. Balfour,7
the Chair of Botany at Edinburgh, and working with several scientific organizations, in particular the
Botanical Society of Edinburgh and the Royal Society of Edinburgh. At the time of his leaving
Edinburgh, Lawson’s work in botany and microscopy was well known and documented in over 50
published articles.7
At Edinburgh, the first Chair of Chemistry, the second in Britain after Cambridge, was
appointed in 1713, joining botany (1676)10 in the Medical Faculty.11 From this early beginning,
Edinburgh maintained a strong position in chemistry, with such scientists as Joseph Black,
discoverer of carbon dioxide, and Alexander Crum Brown, developer of graphical chemical formulae,
on the faculty at various times.11 Lawson certainly acquired his knowledge of chemistry during his
time at Edinburgh.
Page 4 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
5/27
It is recorded that Lawson was awarded the D. Phil. degree in 1857 from the University of
Giessen.12 This record also indicates that there was no thesis associated with his degree. A more
detailed listing of doctoral degrees with theses and promoters contains no mention of Lawson.13 In
fact, a listing of personnel at Giessen in 1856–57, “Personal-Bestand der Großherzoglich
Hessischen Ludewigs-Universität Giessen”, from the archives shows no record of Lawson. Thus, it
is doubtful whether he ever studied in Germany,9 and his D. Phil. degree must have been awarded
in absentia. This was not uncommon; the majority of the degrees from Giessen in 185712 were
awarded without submission of a thesis. There is also no record of what published work Lawson
submitted as a basis for the awarding of his degree. However, the possession of a degree from
Giessen would have helped him to obtain an appointment in chemistry, since at that time, Giessen
was acknowledged to have state-of-the-art facilities in chemistry, largely due to the work of Justus
von Liebig, who was widely recognized for his teaching and broad-based research in chemistry.
In 1858 Lawson accepted an appointment as Professor of Chemistry and Natural History at
Queen's College in Kingston, Upper Canada. Over five years in Kingston, he taught chemistry and
set up laboratories in chemistry and botany. Many students were influenced to take up botany
through contact with Prof. Lawson,8 who was also instrumental in establishing the Botanical Society
of Canada.14 Indeed, he is considered by many to be the “Father of Botany in Canada”.
From 1863 until his death in 1895, Lawson taught chemistry in both the Faculty of Arts and
the Medical Faculty at Dalhousie. A practical man, he became involved with agriculture. He not only
owned and ran a model farm a short distance from Halifax, but also acted as secretary to the Central
Board of Agriculture of Nova Scotia and subsequently as Secretary of Agriculture for the Province.
He was also a member of many learned societies, including the Royal Society of Canada, of which
he was a charter member and president in 1887–88,7 and the Nova Scotian Institute of Science, of
which he was president at his death.8 In addition to these teaching and other duties, he continued to
publish scientific research, mainly in botany. All told he authored over one hundred publications,
Page 5 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
6/27
which included chemical topics such as nomenclature, water analysis, chemistry and heat, gold
processing and the discovery of a dye from the black spruce Aphis.6
Changing chemical concepts in the eighteen sixties
To provide some perspective for the chemistry in Lawson’s lectures, which preceded the
famous periodic table published by Dmitri Mendeleev in 1869,15 it might be helpful to list a few
developments in the state of chemical knowledge around that time. In 1852, Edward Frankland is
credited with proposing the theory of chemical valency,16 and the tetravalency of carbon was
suggested by August Kekulé in 1857–58.17 In the 1860s and beyond, the Theory of Structure, which
recognized and exploited specific atomic connectivity, was developing,18 and the ring and bonding
structure of benzene was suggested by Kekulé in 1865.19 In the mid-1850s Kekulé and other notable
German chemists published using the old atomic mass scale championed by Berzelius in which C =
6 and O = 8.17 At the Karlsruhe Congress of 1860, held as an attempt to bring some order and
consistency to the state of chemistry at that time, Stanislao Cannizzaro used Amedeo Avogadro’s
ideas about diatomic elemental gases and gas volumes to propose a consistent set of atomic
masses, which included C = 12 and O = 16.20,21 Although not immediately accepted, these masses
eventually became universally employed.
Chemistry as taught by Lawson in 1868‒69
A detailed account of the 89 lectures presented daily by Prof. Lawson was recorded by
Alexander Russell,22 one of nineteen BA students in the chemistry class who later became a Minister
in the Presbyterian Church. Russell’s notebook,23 located in the Dalhousie University Archives,
describes a remarkably comprehensive presentation of chemistry, with frequent lecture
demonstrations. The lecture notes covered 185 pages of fine handwriting. As in any note-taking
situation, it is possible that the notes may include some misrepresentations of the material
Page 6 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
7/27
presented. Also, it must be noted that some of the many misunderstandings prevailing at that time,
in particular, the problem of the atomic masses, explained below, were reflected in the recorded
material. The recommend text for junior students was “Practical Chemistry” by Stevenson Macadam
(W. & R. Chambers, Edinburgh); for senior students, “Fownes Chemistry for Students” by George
Fownes, edited by Robert Bridges (Blanchard & Lea, Philadelphia) was recommended. What
follows is a selected summary of the large amount of material detailed in Russell's notes.
Names and Formulae
The elements were described using the same symbols used today but atomic masses were
called “equivalent numbers”. While some of the atomic masses were very close to those accepted
today, a considerable number were half of current values (e.g. C = 6 and O = 8) and a few were
inexplicably different (see Table 1). This led to differences in chemical formulae compared to present
practice. If a compound contained only elements with correct atomic masses or masses incorrect by
a factor of two, the chemical formula would agree with the contemporary formula. If a compound
contained a mixture of these types of elements, the number of atoms of some elements in the
formula would be wrong by a factor of two. For example, Lawson presented the formula of water as
HO, not H2O. The formula HO gave the correct weight percent of the elements in water assuming
atomic masses of H = 1 and O = 8.
In addition to the classification of elements as metals or nonmetals, some elements with
similar properties were grouped, such as Na, K, Mg, Ca, the “alkali class”. Compounds such as
oxides and sulphides were recognized, and formulae were written as today using subscripts to
indicate stoichiometry. Basic oxides were called “oxides”, while acidic ones were called “acids”, but
not all acids contained oxygen, e.g. HCl. When writing formulae for binary compounds, the metal
symbol was written before the nonmetal, whereas in oxides, the basic oxide was first and the acidic
one second. For example calcium sulphate could be written either as CaSO4 or as CaO,SO3, but
potassium sulphate was written as KSO4 or KO,SO3. These formulae reflect the correct mass for
Page 7 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
8/27
potassium and the half values of the masses for calcium, oxygen and sulphur. Formulae for more
complex molecules (called “bodies” by Lawson) were known, such as morphine, C34H19NO6; when
corrected for the erroneous atomic weights of carbon and oxygen, the formula is C17H19NO3, as is
accepted today.
Stoichiometry, thermochemistry and electrochemistry
Using combustion as an example, Lawson introduced fundamental principles. In particular,
he told the students that matter could not be created or destroyed and that one type of matter
reacted with another in fixed proportions. For example, he explained how 1 pound of carbon would
react with 2⅔ pounds of oxygen to produce 3⅔ pounds of carbon dioxide, plus enough heat to
convert 122 pounds of water at 60º F to steam at 212º F, which could be used to do work. The
stoichiometry of the reaction is in agreement with values used today, while the thermochemistry is
off by less than 6%. There is no record in the notes as to how the numbers were derived, but
presumably they were experimental values from other work. While the units used by Lawson are
antiquated, he realized that there were two temperature scales in use, Centigrade (now Celsius) and
Fahrenheit, and that measurement of temperature depended on the expansion of materials, such as
a metal rod or a column of mercury.
After noting that combustion of coal in air produced carbon dioxide, and combustion of
carbon compounds containing hydrogen also produced water (shown as HO), he pointed out that all
living animals respire, consuming oxygen and producing carbon dioxide. To maintain balance on the
planet, he noted that plants absorb carbon dioxide, release oxygen and convert carbon dioxide into
coal over a long period of time, thus storing energy for later use. It is sobering to realize that this was
taught to students 150 years ago and is of equal importance in the world today; sadly however, this
is unrecognized by many.
Lawson also indicated that combustion produced both heat and light. Within this context, he
pointed out how various forces could be converted into other forces, such as steam generation
Page 8 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
9/27
(above), or the use of electricity to overcome chemical affinity in the electroplating of copper from a
solution of cupric sulphate. Other topics discussed included batteries made from zinc and copper,
and the flow of electricity in a coil used to produce a magnetic field.
Properties and reactivity of common nonmetallic elements
Oxygen
After briefly mentioning the discovery of oxygen by Priestley, Lawson pointed out that it could
be obtained by applying a force (i.e. heat) to many oxygen-containing substances, although it was
usual to do this with chlorate of potassium, which he wrote as KO,ClO5. Lawson demonstrated the
preparation of oxygen and did the usual experiment in which a glowing wood splint placed in an
atmosphere of oxygen would burst into flame. He warned that saltpetre (potassium nitrate) could
behave similarly, but was dangerous on contact with combustible materials and he classified
materials as either combustible or not. He suggested that oxygen had powerful attraction for other
elements, reacting with all except fluorine, although that perhaps it did, but this had not been
discovered. Fluorine was dangerous to handle.
Lawson explained how certain metals could react with oxygen in the air, but the oxides
formed then inhibited further reaction; this was not true for iron, which did react further. Oxides could
be basic, neutral or acidic, and the more oxygen the latter possessed, the more acidic they became.
Finally, Lawson discussed the dilemma of oxygen and ozone. It was known that these were
different, but the difference was ascribed to them being allotropes, such as red and white
phosphorus. Ozone was formed after lightning strikes, was a powerful oxidant and, like chlorine, an
excellent disinfectant and bleach. Ozone was easily detected by starch paper soaked with iodide of
potash and the depth of the resulting blue colour indicated the concentration.
Hydrogen
Page 9 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
10/27
In a discussion of hydrogen, the lightest element with an equivalent number of 1, Lawson
noted that it combined with other elements in the proportion of the relative equivalent numbers.
Hydrogen and oxygen could be generated at the different poles in the system by the electrolysis of
water. A different way was to add a small piece of sodium metal to water. He also pointed out that
other less reactive metals, such as zinc, would do the same when placed in a dilute acid and he
demonstrated both effects. He warned that a mixture of oxygen and hydrogen was perfectly stable if
undisturbed, but would explode if subjected to any forces such as an electric spark, light, sudden
compression or finely divided platinum. In another of many in-class demonstrations, Lawson
exploded a mixture of hydrogen and oxygen. Hydrogen and oxygen could be ignited in a controlled
jet to produce an intensely hot flame. If the flame was confined in a hollow glass tube, sound would
be generated at a frequency which could be tuned by the size of the tube. Lawson explained this as
the flame being a continuous series of small explosions, thus generating sound. Burning of hydrogen
produced its oxide, water. Hydrogen was the lightest known gas with 100 cubic inches weighing only
2.14 grains and in principle it could float on top of other gases, but in fact mixed with them because
of diffusion. He mentioned that hydrogen had a second oxide, HO2, which was an excellent
bleaching agent. Hydrogen could form an oxide which was both acidic or basic.
Nitrogen
Lawson told the class that the properties of nitrogen were best characterized by negatives –
it did not support either life or combustion, was colourless, tasteless, non-toxic and insoluble in
water. He divided the components of air into major, nitrogen and oxygen, and minor, carbon dioxide,
water vapour, hydrogen sulphide, nitric acid, ammonia, ozone and sulphur dioxide. Nitrogen had a
weak affinity for oxygen, but it did form five oxides, each with an increasing proportion of oxygen.
These were given as NO (nitrous oxide, laughing gas), NO2 (nitric oxide), NO3 (hyponitrous acid),
NO4 (nitrous acid) and NO5 (nitric acid); these formulae have a correct atomic mass for nitrogen, but
an incorrect one for oxygen.
Page 10 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
11/27
Heating nitrate of ammonia to 300º F in a retort gave water and nitrous oxide, because
hydrogen had a greater affinity for oxygen than did nitrogen. Humphrey Davy had found that the gas
could cause a form of intoxication, hence the name laughing gas. Nitrous oxide acted on the human
muscular system, not the nervous system, unlike the actions of morphia or strychnine. Nitric oxide
was prepared by heating nitrate of copper and was a gas with a strong disagreeable odour. It could
be oxidized to NO3 or NO4, which were unimportant.
Nitric acid was a strong oxidizing agent and could be formed by heating 2 parts of nitrate of
soda with one part of sulphuric acid; nitrate salts were common in the deserts of South America.
Nitrogen and oxygen do not combine directly, only in an electrical discharge. In large cities with little
pure air and much filth, nitric acid can even be found on the walls of buildings. Nitrates are formed
by decaying organisms and have been found in wells close to churchyards, a point relevant to
environmental concerns today. Nitric acid was normally obtained as a 70% aqueous solution, but it
can also be prepared in essentially anhydrous form. It is colourless when pure, but is brown when
contaminated with nitrogen dioxide. It reacts vigorously with most metals and with living or dead
tissues.
Carbon
Carbon was interesting as it is found in several forms, wood charcoal, lamp black, animal
charcoal, coke, graphite (soft and considered to be the normal form of carbon) and diamonds (the
purest form of carbon and very hard). Animal charcoal had the greatest porosity and was used to
purify sugar in its manufacture. Lawson presented a table showing the relative amounts of gases
that could be adsorbed by animal charcoal, with ammonia being the greatest and hydrogen being
the least.
Carbon was known to form two oxides, CO (carbonic oxide) and CO2 (carbonic acid). The
former was formed during incomplete combustion of charcoal and was poisonous to the lungs, but it
could also be made by treating ferro cyanide of potassium with sulphuric acid. Lawson explained
Page 11 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
12/27
that this reaction was complex, but there was only enough oxygen present to form CO and not CO2.
The latter was a heavy gas, formerly known as “fixed air” because it could be fixed as solid calcium
carbonate when bubbled into lime water. It could be released from this, or from marble, by heating or
by treating the solid with acids such as sulphuric or hydrochloric. It was freely soluble in water to
form the weakest known acid. Carbon dioxide was also produced by fermentation, respiration,
volcanic eruptions and in mineral springs.
Lawson moved on to discuss compounds of carbon with hydrogen, while pointing out that
most of these would be discussed as organic chemistry. He focused on the two simplest
hydrocarbons, CH2 (light carburetted hydrogen) and C4H4 (olefiant gas). The former had a long
history of use for illumination and was made commercially by dry distillation of bituminous coal, In a
detailed discussion he described how impurities such as ammonia, hydrogen sulphide, sulphur
trioxide and carbon dioxide were removed and analyzed to ensure that the coal gas was pure.
Lawson explained how light came from combustion of carbon and not from combustion of hydrogen.
The halogens
The elements chlorine, bromine, iodine and fluorine form a natural grouping. In nature none
is found as a free element, but always as a salt, the first three in salt deposits and sea water, the last
in the mineral fluorspar. Of these, chlorine was covered most extensively initially by pointing out that
the action of a strong acid on the muriate of soda produced hydrogen chloride. To make chlorine,
the hydrogen needed to be removed, which was done by adding manganese dioxide. Chlorine had a
great affinity for all elements and was dangerous to handle; it attacked the lungs causing violent
bursts of coughing and usually permanent damage. Chlorine could be liquefied under pressure, but
it did not congeal even at the coldest temperatures. Chlorine reacted unusually with some organic
compounds such as turpentine; one hydrogen atom was removed and replaced by a chlorine atom.
We now know that this is an example of free radical halogenation at an allylic centre. Chlorine was
soluble in water producing unstable hypochlorous acid, but the sodium salt of this acid was stable
Page 12 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
13/27
and, like chlorine, was an excellent bleach and disinfectant. Lawson described four oxides of
chlorine, as well as their associated acids, each with an increasing oxygen content.
Hydrogen chloride was of great importance, four volumes of which were formed explosively
when two volumes each of chlorine and hydrogen were exposed to sunlight. It had a great attraction
for water, which could absorb 480 times its bulk of hydrogen chloride. Hydrochloric acid would
dissolve most metals except gold and platinum, but if mixed with nitric acid to make aqua regia, even
these could be dissolved. It was chlorine that was set free in this acid which attacked these noble
metals.
Bromine could be released from its sodium salt by the action of chlorine and was easily
separated because of its solubility in “sulphuric ether” (ethyl ether). Where chlorine was a green gas,
bromine was a brown liquid and iodine a violet solid. In general bromides were much more soluble in
water than chlorides. Iodine was unusual because on heating the solid converted directly to the
vapour. It was found in many marine organisms, especially kelp.
Fluorine was different from the other halogens. Only hydrogen fluoride was discussed.
Lawson prepared this from calcium fluoride and sulphuric acid to demonstrate how it etched glass.
He emphasized that hydrogen fluoride was dangerous and caused severe burns to human tissues.
Sulphur, phosphorus, silicon and boron
Sulphur was discussed at length, initially by describing its allotropes, crystalline, vitreous
(glassy) and amorphous. The first was soluble in bisulphide of carbon, the second was “elastic” and
the third was “earthy”. Sulphur was found in volcanoes, pyrites and in all living matter, especially
albuminous matter such as eggs. Lawson described seven oxides of sulphur, with emphasis on the
dioxide, which was easily oxidized to the trioxide. Sulphur dioxide and its aqueous solution were
good bleaching agents, but differed from the bleaching action of chlorine. Sulphuric acid was the
strongest known acid, formed salts with alkalis and oxidized metals to form salts, as Zn + HO,SO3
gave ZnO,SO3. Hydrogen sulphide, written as HS, was found in decaying vegetable or animal matter
Page 13 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
14/27
and was discussed in detail. It was used as a test for metals, with many of them giving insoluble,
coloured sulphides. Lawson made only a brief mention of selenium, pointing out that it was always
found in conjunction with sulphur and that “HSe smelled even worse than HS”.
Phosphorus never occurred in nature in the free elemental form, but was always combined
with oxygen, although occasionally it was found as a phosphide. Phosphorus mostly occurred as
phosphate of lime, 3CaO,PO5, which was found in soils and bones. The element was essential for
life and even occurred in the brain. Lawson described the white and red allotropes of phosphorus.
The former was toxic, smelly and soluble in bisulphide of carbon; it burned in air and glowed in the
dark. This latter property was a sensitive test for the material. By contrast, red phosphorus was
more stable and much less toxic or soluble. The element burned with a brilliant flame and emitted
copious quantities of white smoke. It was used in matches. Lawson described four oxides of
phosphorus including the most common, “PO5, phosphoric acid”. Next to nitrates, soluble
phosphates were the most important components of manures for growing plants.
Finally Lawson briefly discussed silicon and boron. The former occurred mainly as the
dioxide, the only known oxide. Glass was a silicate of potassium. Silicon also formed a compound
with fluorine, which he wrote as “SiF”. Boron was somewhat like silicon and was found in nature as
borates (salts of boracic acid). These could be found in volcanic springs in places like Persia.
Borates have been used to render metals fusible or as “a bromide flux”. These two elements
completed coverage of the nonmetals; the lecture on December 23, 1868 was the last of the
calendar year.
Properties and reactivity of selected metallic elements
Lawson spent the first four weeks of the new year discussing a total of thirty-two metallic
elements. He first covered the alkali and alkaline earth metals as groups, followed by most of the
first row of transition metals in reverse order to their atomic masses, then a mixture of elements in
no particular order, and finally some of the noble metals. If judged by the volume of notes, Lawson
Page 14 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
15/27
may have considered Au, K, Al, and As to be the most important metals and Ti, V, Li, Mo, W, and Pd
to be the least important. It is more likely that these two groups represented the metals about which
the most and least were known. All this was presented in a descriptive manner and included such
topics as properties, methods of preparation, uses, chemical reactions, and analytical tests. Four
examples of elements important to Nova Scotia are included below.
Sodium
Lawson noted that sodium metal behaved in a manner similar to potassium. It was soft,
lustrous when first cut, soon reacted with air and burned with a yellow flame when placed in water,
resulting in an alkaline solution. Two oxides of sodium were mentioned: NaO, caustic soda ‒ by far
the more important of the two, and Na2O. Various methods of preparing carbonate, bicarbonate
(important in baking), and sulphate of soda were described. Originally sodium salts were obtained
from plant ash but more recently sea salt (NaCl) had replaced this source.
Iron
Lawson described iron as a very important element with a wide distribution in various rocks
and soils as well as in our bodies, especially in blood. Iron ores included both FeO and Fe2O3, “bog
ore”, and iron pyrite (FeS2). He suggested that magnetic iron ore (magnetite) was a mixture of FeO
and Fe2O3. Iron could be obtained by roasting some of the iron ores in air to form iron oxide. The
oxide was then placed in a blast furnace along with coal and a stream of air passed through the
furnace, a structure described as 50 feet in height and 17 feet in diameter. The coal provided heat
and took the oxygen away from the iron, providing elemental liquid iron, which was removed from
the furnace.
Since iron was such an important element, tests were needed for its detection. Two tests
were described. The first involved adding the ferro cyanide of potassium to an iron-containing
solution, which resulted in a blue colour and a dark blue precipitate. The second used sulfo cyanide
Page 15 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
16/27
of potassium (potassium thiocyanate) as a reagent to form a deep red solution indicating the
presence of iron. This was described as a very sensitive test, useful for mineral waters. Other
aspects of iron chemistry were to be discussed in later sections of the course, where appropriate.
Arsenic
This element had been known for a long time. Arsenic shared some properties with
phosphorus, was quite poisonous but otherwise had a high metallic lustre and electrical conductivity.
It was said to burn in air with a bluish flame and its vapours smelled like garlic. The main source of
the element was mispickel, a mixture of As, Fe, and S (today known as arsenopyrite). Other sources
included the arsenides of Ni or Co. These ores could be roasted in air in a “reverberating furnace” to
give AsO3. Solid arsenic was collected on the walls of long, cool tubes associated with the roasting.
Two oxides of arsenic were known: AsO3 (arsenous acid), the more important, and AsO5
(arsenic acid). Arsenous acid was not very soluble, even in boiling water. Thus, even though arsenic
was recognized as poisonous, if this property were of interest, strychnine was recommended over
AsO3 for solubility reasons. To prepare soluble arsenic, salts like KO,AsO3 should be employed –
this was the form of arsenic used in medicine. A mixture of As and Cu (probably copper arsenite)
produced a green pigment used in paints and in colouring women’s dresses.
An interesting flame test for arsenic in a solid sample was described. The sample was placed
in a glass tube through which hydrogen was passed, generated by adding zinc to sulphuric acid.
This was lit at the end of the tube. The sample was then heated with a “spirit lamp”. If it contained
arsenic the flame changed to a “peach-blossom” colour and white AsO3 was given off. Holding a
porcelain cup near the flame resulted in a coating of solid As, the “Arsenical Mirror”. This is known
as the Marsh test and depends on the formation of arsine in the heated tube. Another test involved
placing a copper foil in a solution containing arsenic. Hydrochloric acid was added to the sample,
which was heated carefully, and a dark coating of arsenic would form on the copper. The copper foil
Page 16 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
17/27
was then transferred to a dry test tube and heated to drive off most of the arsenic. If a red powder
remained, described as an alloy of copper and arsenic, the presence of arsenic was indicated.
Gold
Gold was often found in the elemental form, known by its bright colour, and sometimes found
in combination with As, Sb, Cu, Os, or Ir. In Nova Scotia it was found in quartz veins. Lawson went
into detail about how gold might be formed in such veins. He described the mercury amalgam
method of removing gold from its ores. Solid gold was retrieved from the amalgam by volatilizing the
mercury. He did not mention the health hazards of such a procedure. Gold was characterized as a
yellow coloured metal that never tarnishes, extremely ductile and malleable, and that melts at 2076º
F (contemporary value = 1947º F). It was a good conductor of heat and electricity.
Two oxides of gold were mentioned, AuO and AuO3. It was also noted that gold could form
chloride, bromide, and iodide salts. It did not dissolve in acids except for aqua regia. This formed the
basis of the “acid test”, in which a drop of nitric acid (NO5) was added to the metal. If it dissolved, it
was not gold. Another test involved adding sulphate of iron (FeO,SO3 ‒ ferrous sulphate) to a
solution. If gold was present, a metallic precipitate would form. Other tests for gold in nature were
described. Iron pyrite (“fool’s gold”) could be differentiated from gold using a hammer – the iron
pyrite would fracture, gold would not. Or a drop of nitric acid would form fumes with iron pyrite but
not with gold. Copper pyrites could also appear to be gold – these fused with heat while gold did not
(unless quite hot). Finally Lawson mentioned that mica could be mistaken for gold due to its glitter in
rocks but did not propose a test to determine its presence.
Mineralogy
Five lectures were given to the discussion of mineralogy. Lawson believed that the
importance of mineralogy had waned somewhat over the fifty-year period preceding his lectures. He
Page 17 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
18/27
made a distinction between minerals, which have a fixed composition, and rocks, which consist of
mixtures of minerals. He suggested that geologists were more interested in rocks than minerals.
Minerals were classified into seven orders. Each order consisted of various families. Lawson
described the properties of the major minerals in the more important mineral families. These
properties consisted mainly of cleavage and fracture, hardness, colour and lustre, and chemical
composition. He recognized that the colour of a mineral often depended on the impurities present in
the mineral. As an example, orthoclase belonged to the felspar (Lawson’s spelling) family of order
one, oxidized stones. It was said to be a double silicate of aluminum and potassium. Lawson did not
list the exact composition, known today to be KAlSi3O8, but indicated that it consisted of alumina,
silica, and potassium oxide. The other properties consisted of specific gravity, 2.5; hardness, 6,
(somewhat softer than quartz which had a hardness of 7 to 7.5 on the scale used by Lawson); and a
colour ranging from colourless to white, but light yellow or pinkish due to impurities.
The material presented on mineralogy was entirely descriptive.
Organic and biological chemistry
Lawson devoted almost one third of his lectures to this theme, paying particular attention to
biological concepts, not surprisingly since most organic compounds known at that time had a
biological origin. Since he was teaching medical students, he devoted time to anatomical concepts,
as well as to therapeutic uses and toxicity of compounds. He distinguished between tissues and
products. Mostly the elements involved were C, H, N, O, P and S.
His introductory remarks were especially interesting: “The object of the chemist is to bring
together the elements. Each of the elements we find has a certain amount of ‘force’. The chemist
has to bring them together and they will form a new body. In order for particles to combine they must
be in contact, e.g. if metals are mixed and heated they will form alloys. Bodies can be broken up by
disturbing the chemical affinity, by application of some other force, such as heat, light, electricity or
mechanical action. Each element then takes back the chemical affinity that it originally possessed. In
Page 18 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
19/27
each case, chemists are only arranging bodies, making possible combinations, not by direct force
not by the Alchemists Philosopher's Stone, but simply by placing them together. In inorganic
chemistry we always know what will come, what can be done in nature can always be done by the
chemist. In organic chemistry the same compounds are never formed from their elements, they are
formed differently. Compounds are formed because there is something more powerful than chemical
affinity, maybe the vital force.”
Clearly at that time, the concepts of chemical structure were in flux and Lawson was
obviously unsure of how organic compounds and structure were related. His suggestion of “vital
force” appears to be in relation to the fact that plants and animals can make different organic
compounds and today we know that this is mostly done by enzyme-catalyzed reactions. He
proceeded to give an example of how plants could take carbon dioxide and water from the air and by
“some force” could split the carbon dioxide to give oxygen with “CHO” left behind, but this was not
further deoxygenated. Different plants had different powers of making secretions, e.g. balsam firs
made oils containing only carbon and hydrogen, but plants could also make compounds consisting
of carbon, hydrogen and oxygen.
Lawson described the general anatomical details of both plants and animals. Plants were
composed of cells which could be seen under a microscope. Cells could contain sugar, starch and
lignin, while leaves contained chlorophyll. Cells have a nucleus which enables new cells to be
formed. Different tissues could be classed as cellulosic or woody, whereas animals had different
classes of tissues. Alveolar tissues formed the basis of animal structure, adipose tissues contained
fats, cartilaginous tissues contained hyaline substance, while osseus tissues were filled with
CaO,CO2 and 3CaO,PO5. All tissues were permeated with liquids.
The saccharine group
After this general introduction, Lawson moved on to discuss the saccharine group, including
starch, lignin, cane sugar, fruit sugar and milk sugar, along with their empirical formulae. Starch,
Page 19 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
20/27
found only in plants, was important and he gave details of its manufacture and detection by chemical
tests. He pointed out that pure sugar did not ferment, but in the presence of yeasts or fungi, sugars
would ferment to alcohol, carbon dioxide and water; the alcohol could be obtained by distillation.
This provided an entry to the homologous series of alcohols, which concluded with “melisic alcohol,
C60H62O2”, today known as 1-triacontanol. The next concepts defined the same homologous series
of radicals, written as CnHn+1, beginning with methyl, then ethyl, followed by derivatives of these
such as alcohols, aldehydes and acids. Lawson introduced the radical formyl as C2H, which he
pointed out could be considered as derived from methyl by loss of hydrogen. The simplest acid,
formic acid, was found in red ants. Related to this, he mentioned chloroform, which produced a state
of insensitivity to pain. Ether does this also, but is less powerful.
Discussion continued to “acetyl” and higher homologs; acetic acid was written as “C4H3O3 +
HO”. Oxalic acid could be formed by oxidizing “amyl” or fermenting potato or grain starch. One
hundred pounds of starch could be fermented from which 150 pounds of crystalline oxalic acid could
be obtained. This acid was toxic with the antidote being lime or chalk which formed an insoluble salt.
Oxalic acid was good for removing ink stains, it did not grow mould. Lawson also discussed other
plant-derived compounds, including the conversion of wood fibres into paper.
Fats, oils and organic acids
A description of fats and oils, such as stearine, oleine and palmitine, led to drying oils (poppy,
linseed, cod liver, hemp, etc.), non-drying oils (olive, almond, rape, animal, etc.) and intermediate
oils such as ricine (castor oil). Most oils had an even number of carbon atoms, although margarine
yielded margaric acid which had an odd number. The drying oils absorbed oxygen from the air to
make resins, with linseed oil being the best at this. By contrast lubricating oils were unreactive and
did not absorb oxygen.
Fats could be saponified by heating with alkali such as sodium or potassium hydroxide to
give the salt of the fatty acid. Potassium salts were more soluble and made soft soaps, whereas
Page 20 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
21/27
sodium salts gave hard soaps and calcium salts were insoluble in water. The topic of oils was
concluded by discussing volatile oils, such as balsam, camphor, cedar, coriander, ginger and lemon;
all were plant-derived.
Lawson moved on to discuss polybasic organic acids, those which combined with two or
more equivalents of base; among others, tartaric, malic and citric acids were featured. He described
a range of derivatives of tartaric acid, including among others, cream of tartar, tartar emetic and
Rochelle salt. Lactic acid was obtained from fermenting milk, which contained lactose and casein,
the latter being important for building muscle. Indeed, milk was the perfect food. Meconic acid was
found in the milky juice of certain poppies, from which opium was obtained. Tannic acid was found in
the bark of certain trees and was used for tanning leather, after which Lawson discussed glues and
inks.
Nitrogen-containing organics
Alkaloids were described as nitrogen-containing, strongly basic poisons found in plants.
They could be classified as either volatile oils, or crystalline solids. Examples were given of both
types. Nicotine, a component of tobacco was important and for customs and excise purposes,
tobacco leaves could be identified by treating with boiling water and smelling the extract. Lawson
described the extraction of nicotine from tobacco leaves and pointed out the toxicity of the alkaloid;
even a small quantity was able to kill a large animal. He discussed some fifteen alkaloids in five
lectures, concluding with caffeine and theobromine. He then discussed organic dyes followed by the
chemistry of wood and its derivatives such as gun cotton.
Coverage then turned to other nitrogen-containing organics, the first being cyanogen, which
Lawson also called “bicarbide of nitrogen”, a misnomer, since in chemistry today carbides have
carbon combined with a less-electronegative element. He explained that cyanogen was important in
providing support for the theory of radicals, although the theory of radicals was supplanted some ten
years earlier by the theory of valency propounded by Frankland and Kekulé. Heating cyanogen with
Page 21 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
22/27
potassium gave potassium cyanide, one of a range of compounds involving “cyanide”, such as
potassium ferricyanide and Prussian Blue. Lawson also described the isolation and properties of
hydrocyanic acid, a weak acid which just turned blue litmus red and was the most toxic substance
known.
Medical connections
Lawson did not mention the word “biological”, but devoted the last eight lectures of the class
to biological and medical aspects of animals and how these depended on plants for survival. He
pointed out that animal matter gave a strong smell of ammonia when burned, but plant matter did
not. Albuminoids included proteins, albumin, fibrin (muscle tissue), casein and legumin and were
plastic materials. Gelatinous substances included gelatin, chondrin, osseine, sclerotin and cartilage.
He described functions for most of these materials along with properties, such as pure albumin could
be obtained from egg white and was precipitated by the addition of acid. Lawson presented a table
classifying CHO-containing organic molecules into seven groups. All of these, whether animal or
plant-derived, came from carbon dioxide and water, but also with a deoxygenation step that only
plants could do. For example, oxalic acid C4H2O8, was derived from 4CO2 + 2HO, but two oxygen
atoms needed to be lost.
He reiterated how animals get food from plants, which are the “servants of animals” while
plants generate “respiratory food” by means of the sun and air. He discussed how food needs to be
well masticated by saliva and then continued to discuss digestion, the stomach, intestines and
pancreatic fluid. Tape worms absorb nutrients, but have no stomach. The final topics were urea and
cyanic acid together with blood, which contained iron. Identification of blood stains was normally
done using a microscope.
Final Examinations
Page 22 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
23/27
The final examinations for all Dalhousie courses were printed in the Calendar of the following
year.2 Students in Professor Lawson's class of 1868−69 had access to the junior and senior
examinations in the previous year. A period of four hours was provided for these written
examinations, which consisted of eight to twelve questions. In general, one question on each
examination involved a mathematical calculation.
From these examinations it is obvious that the junior and senior students were responsible
for different parts of the material. Russell and his fellow students were expected to have memorized
the equivalent numbers of all the common elements as well as all the descriptive chemistry that
Lawson had covered. The Junior Examinations mainly covered inorganic topics and included
questions on nomenclature, definitions, chemical formulae, as well as descriptions of chemical
processes and preparations. For example, a question on the Junior Chemistry Examination of 1868
reads “Describe fully the Chemical changes by which Common Salt is converted into Carbonate of
Soda, in the ordinary methods of the Soda Manufacture.” In 1869, Russell was asked about silver in
solution, how and in what form it could be precipitated. The final part of this question was to provide
a calculation as to how one could estimate the amount of metallic silver.
Senior Examinations emphasized organic and biological chemistry with greater weight
placed on description of more complex processes. The Senior Examination in 1868 required a
calculation as follows: “Found grains 3.05 Tersulphide of Antimony; suppose the whole of the Sb to
have originally existed in the form of KO,SbO3, C4H4O10 + 2 HO, how much Tartar Emetic would the
above 3.05 gr. represent?”
Concluding remarks
The remarkable feature of Lawson's lectures is the level of detail that students were
expected to learn. We were surprised by the extensive level of knowledge at that time. However,
not all the material was fully up to date. Keeping abreast of new developments was undoubtedly
Page 23 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
24/27
challenging since few chemical journals were available at that time and the availability of such
material in Halifax is unknown. During the lectures, Lawson mentioned other texts by Wilson
(Chemistry) and Graham (Inorganic Chemistry). Students could sign up for an extra course in
practical chemistry (3rd year) which used C. R. Fresenius' classic text on Qualitative and Quantitative
Analysis (Vieweg, Braunschweig). It was clear that there was confusion in understanding some
chemical concepts. Very little was said about chemicals from coal while petroleum-based chemicals
were yet to be developed. The lectures were striking in how Lawson presented the natural balance
between plants and animals and how virtually all of the organic chemicals were plant-derived,
concepts to which much attention is now being focused, 150 years later.
Acknowledgements
We thank Janelle S. Ramaley (Universität Heidelberg) and Lutz Trautmann, (University
Archives, Justus-Liebig-Universität Giessen) for their help in obtaining information regarding the
awarding of doctoral degrees at Giessen in the middle of the nineteenth century. We also thank the
staff of the Dalhousie University Archives for their assistance.
References (all websites referenced below were accessed in November, 2017)
(1) Chute, W. J. “Chemistry at Dalhousie”, 2nd ed., Dalhousie University, Halifax, NS, 1986.
(2) https://dalspace.library.dal.ca/handle/10222/11502. (Dalhousie University Calendars).
(3) https://books.google.fr/books?id=i-
INAAAAQAAJ&printsec=frontcover&hl=fr&source=gbs_ge_summary_r&cad=0#v=onepage&q&f=fals
e. (McGill University Calendar, see page 33).
(4) https://archive.org/stream/qucalendar_1867#page/n19/mode/2up. (Queen's University Calendar).
(5) https://archive.org/stream/uoftcalendar1867#page/26/mode/2up. (University College, Toronto
Calendar).
Page 24 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
25/27
(6) Anon. Bibliographies of Members of the Royal Society of Canada, Proc. Royal. Soc. Can. 1895,
12, 49-52.
(7) Zeller, S. “LAWSON, GEORGE” in Dictionary of Canadian Biography, Vol. 12, University of
Toronto/Université Laval, 1990. http://www.biographi.ca/en/bio/lawson_george_12E.html.
(8) MacGregor, J. G. Proc. Nova Scotian Inst. Sci. 1895, 9, xxiv – xxx.
(9) Rousseau, J. and Dore, W. “L’Oublié de L’Histoire de la Science Canadienne – George
Lawson, 1827–1895” in Pioneers of Canadian Science, ed. G.F.G. Stanley, Toronto, 1966, pages
54-66.
(10) The University of Edinburgh, “Our History, Botany”.
http://ourhistory.is.ed.ac.uk/index.php/Botany.
(11) The University of Edinburgh, School of Chemistry, “300 Years of Achievement”.
http://www.chem.ed.ac.uk/about-us/tercentenary/300-years-achievement.
(12) Kössler, F. “Verzeichnis der Doktorpromotionen an der Universität Giessen von 1801-1884” in
Berichte und Arbiten aus der Universitätsbibliothek Giessen 17, 1970.
http://geb.uni-giessen.de/geb/volltexte/2006/3594/pdf/Koessler-1970-Promotionen.pdf.
(13) Kössler, F. “Katalog der Dissertationen und Habilitationsschriften der Universität Giessen von
1801-1884” in Berichte und Arbiten aus der Universitätsbibliothek Giessen 22, 1971. http://geb.uni-
giessen.de/geb/volltexte/2006/3610/pdf/BA-22.pdf.
(14) Connor, J. T. H. Scientia Canadensis 1986, 10, 3. doi: 10.7202/800223ar.
(15) Royal Society of Chemistry, “History of the Periodic Table”. http://www.rsc.org/periodic-
table/history/about.
(16) Russell, C. A. “The History of Valency”, Humanities Press Inc., New York, 1971, page 34.
(17) Russell, C. A. “The History of Valency”, Humanities Press Inc., New York, 1971, page 62.
(18) Russell, C. A. “The History of Valency”, Humanities Press Inc., New York, 1971, page 142.
(19) Russell, C. A. “The History of Valency”, Humanities Press Inc., New York, 1971, page 242.
Page 25 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
26/27
(20) Everts, S. Chem. Eng. News 2010, 88(36), 60.
https://pubs.acs.org/cen/science/88/8836sci1.html.
(21) Mönnich M. W., Chemistry International, 2010, 32(6), 10-14, originally as Nachrichten aus der
Chemie, 2010, 58 (May), pp. 539–543.
https://www.iupac.org/publications/ci/2010/3206/4_monnich.html#author.
(22) https://findingaids.library.dal.ca/russell-alexander-g-1845-1911.
(23) Alexander Russell's notes from George Lawson's chemistry lectures at Dalhousie College, MS-
2-380, SF Box 40, Folder 16, Dalhousie University Archives, Halifax, Nova Scotia, Canada.
Page 26 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry
Draft
27/27
Table 1. Lawson’s Table of the Elements from Russell’s 1868‒69 Notes Elementa Lawson’s Eq. Atomic Elementa Lawson’s Eq. Atomic
Symbol No.b Massc Symbol No.b Massc
Hydrogen H 1.0 1.0 Arsenic As 75.0 74.9
Lithium Li 6.5 6.9 Seleniumd Se 39.5 79.0
Berylliume Be 6.9 9.0 Bromine Br 80.0 79.9
Boron B 10.9 10.8 Rubidium Rb Unk 85.5
Carbond C 6.0 12.0 Yttrium Y Unk 88.9
Nitrogen N 14.0 14.0 Zirconiume Zr 33.6 91.2
Oxygend O 8.0 16.0 Silver Ag 108.0 107.9
Fluorine F 19.0 19.0 Cadmiumd Cd 56.0 112.4
Sodium Na 23.0 23.0 Tind Sn 58.8 118.7
Magnesiumd Mg 12.0 24.3 Antimony Sb 120.3 121.8
Aluminumd Al 13.7 27.0 Iodine I 127.0 126.9
Silicone Si 21.3 28.1 Cesium Cs Unk 132.9
Phosphorus P 32.0 31.0 Bariumd Ba 68.6 137.3
Sulphurd S 16.0 32.1 Lanthanum La Unk 138.9
Chlorine Cl 35.5 35.5 Cerium Ce Unk 140.1
Potassium K 39.0 39.1 Didymiumf D 48.0 ---
Calciumd Ca 20.0 40.1 Terbium Tb Unk 158.9
Chromiumd Cr 26.7 52.0 Erbium E Unk 167.3
Manganesed Mn 27.6 54.9 Gold Au 197.0 197.0
Irond Fe 28.0 55.9 Mercuryd Hg 100.0 200.6
Cobaltd Co 29.5 58.9 Thallium Thl Unk 204.4
Nickeld Ni 29.6 58.7 Leadd Pb 103.7 207.2
Copperd Cu 31.7 63.5 Thoriume Th 59.6 232.0
Zincd Zn 32.6 65.4 Uranium U Unk 238.0
a Ordered by atomic number
b Lawson’s equivalent to atomic mass
c Present values
d An element with the atomic mass divided by two
e An element with an incorrect atomic mass
f Didymium, actually a mixture of praseodymium and neodymium, was considered an element at that
time.
Page 27 of 28
https://mc06.manuscriptcentral.com/cjc-pubs
Canadian Journal of Chemistry